Atq7 1m1617

QUANTITATIVE DETERMINATION OF TOTAL HARDNESS IN DRINKING WATER BY COMPLEXOMETRIC EDTA TITRATION R.J. TORING1 AND R.J.H. BAGUINON2 1 DEPARTMENT OF CHEMICAL ENGINEERING, COLLEGE OF ENGINEERING 2 DEPARTMENT OF CHEMICAL ENGINEERING, COLLEGE OF ENGINEERING UNIVERSITY OF THE PHILIPPINES, DILIMAN QUEZON CITY, PHILIPPINES DATE PERFORMED: SEPTEMBER 24, INSTRUCTOR’S NAME: MR. AMADO

2015 LUIS ORDEN

1. The concept of complexometric titration and water hardness (why is it expressed as ppm CaCO3?). In a complexometric titration, a metal ion reacts with a suitable ligand to produce a complex. Just like other titration, the equivalence point can then be determined by an indicator. And, water hardness was originally a measure of the water’s ability to precipitate soap. But since calcium and magnesium are the main cause of this, water hardness can now be defined as the amount of calcium and magnesium present which can be expressed into ppm of calcium carbonate. In this experiment, EDTA

(a ligand) was allowed to react with the Ca2+ and Mg2+ ions (metal ions) present in the water during the titration process. The water hardness will then be measured as the ppm of calcium carbonate. [1][2][3] 2. The use of EDTA as complexing agent and titrant. The ethylenediaminetetraacetic acid (EDTA) can react with metal ions in a 1:1 ratio to form very stable complexes at high pH. For this experiment, EDTA would readily react with Ca2+ and Mg2+ in a 1:1 ratio to

form complexes at about pH 10,

making it a good complexing agent and titrant. These behavior is due to its hexadentate structure which allows it to have six donor atoms and thus six sites of attachment for the metal ions, making the formed complexes to be more stable. [1][2][3] 3. The use of EBT as indicator; significance of adding MgCl2⋅6H2O crystals to the titrant. The indicator used in this experiment is the Eriochrome Black T (EBT) gives a good endpoint. EBT forms a colored metal complex at the start which gives the color of wine red. At the equivalence point, the metal ions from the metal complex formed by EBT were removed by EDTA through chelation, making the color turn into blue. This is possible since EDTA binds more strongly to the metal ions than EBT. However, EBT indicator does not react well with Ca 2+ ions because the complex MgIn- is not that strong and wouldn’t give much color. To correct that, MgCl2⋅6H2O crystals were added to give a sharper endpoint. [2][4] 4. The equations titration.

pertinent involved

chemical during

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Initially, EBT reacts with the Mg2+ ions by the equation: Mg2+ + HIn2- ⇆ MgIn- + H+ (blue) (wine red) When EDTA is added, EDTA reacts with the metal ions to form complexes – removing the metal ions in the EBT complex – shown by the equations: Ca2+ + H2Y2- ⇆ CaY2- + 2H+ Ca2+ + MgY2 ⇆ CaY2- + Mg2+ Mg2+ + HIn2- ⇆ MgIn- + H+ MgIn- + H2Y2- ⇆ MgY2- + 2H+ Combining the equations, it would yield the equations that occurs at the endpoint: 2H2Y2- + Ca2+ + Mg2+ → CaY2- + MgY2- + 4H+ 2H2Y + MgIn → MgY2- + HIn2- + H+ 5. The relationship of pH and feasibility of titration using EDTA method for cation analysis; the minimum pH values for cations in EDTA titration. At lower pH, EDTA has its complete 4 protons and exist as H 4Y, decreasing the number of bonding possible to the metal ions. While at higher pH (>10), EDTA is deprotonated and exists as Y4-, giving more sites of attachment for the metal ions. With more sites of attachment, the stable the formed complex will be. The minimum pH values to form a complex for Ca2+ and Mg2+ with EDTA is 7.3 and 10, respectively. [1] 6. Other applications of EDTA complexation. Other uses for EDTA includes industrial, and medicinal and healthsciences uses. For industrial uses, EDTA is helpful in making cosmetic products resistant to air molecules, as purifying agent and preservative,

soften water for more efficient cleansing , preventing the discoloration of dyed fabrics, and reducing reactivity of metal. For medicinal and health-sciences uses, EDTA can be used to treat lead, mercury, chromium, cobalt, zinc (and some other) poisoning through chelation therapy. It can also help to clear bloodstream from unused ions. [5] 7. Difference between experimental and actual value. In the experiment, the sample water isn’t boiled – which makes it possible for some carbon dioxide to dissolve. These carbon dioxide may interfere in the test – giving some difference of the experimental to the actual value. 8. The rationale behind performing the analysis of Ca and Mg at pH 10. As mentioned in the previous answer (number 5), EDTA has more sites for attachment at higher pH, making the complex formed stable. With more stable complexes, the better the result of the experiment is. Also, the minimum required pH for Mg2+ to form a complex with EDTA is at pH 10. Thus, at least pH 10 is required for the complex to form. [1] 9. The effect of using too much buffer in the analysis of Ca and Mg using EDTA titration. With too much buffer, the solution would become too basic. With that, pH changes through titration would hardly affect the solution. This would result to a less sharp endpoint. [1]

10. The correlation between stability of Mg-EDTA complex, CaEDTA complex, MgIn- and CaInand their corresponding Kf values.

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The higher the formation constant of a complex, Kf, the more favored the formation that complex would be – thus, the more stable it becomes. Since Ca-EDTA (Kf =5.0x1010) has higher Kf than Mg-EDTA (Kf =4.9x108), Ca-EDTA is more stable than Mg-EDTA. Also, since MgIn- (K f =1.0x107) has higher Kf than CaIn- (Kf =2.5x105), MgIn- is more stable than CaIn-. Additionally, since the EDTA complexes have relatively higher Kf values than the EBT complexes, thus, the EDTA complexes are more stable than the EBT complexes. [1][4] 11. The rationale behind adding NaOH to the EDTA solution if the sodium-EDTA crystals do not dissolve. The dissolution of sodium-EDTA crystals is normally very slow and is more favorable at higher pH (≥ 8). By adding NaOH, the pH becomes higher and makes the dissolution of sodiumEDTA crystals more favorable. Thus, the crystals would be dissolved faster. [1]

12. Possible sources of errors and their effect on calculated parameters. Possible sources of errors include too much addition of EBT, too much addition of buffer, too fast addition of HCl during CaCO3 dissolution, and overtitration. Too much addition of EBT indicator and buffer can lead to a less sharp endpoint, making inaccurate volume

readings. Too fast addition of HCl during CaCO3 dissolution can make excess formation of CO2 leading to a loss of material. Lastly, overtitration can also make inaccurate data. [1] REFERENCES [1] Skoog, D.A.; West, D.M.; Holler, F.J.; Crouch, S.R. Fundamentals of Analytical Chemistry, 9th ed.; California: Brooks/Cole Cengage Learning, 2014. [2] Winona State University. Determination of the Hardness of Natural Waters. http://course1.winona.edu/mengen/Ch emistry%20320/Lab/Hardness%20of %20Water.pdf (accessed September 30, 2015) [3] The University of Tennessee, Knoxville. Experiment 3: EDTA Determination of Total Water Hardness. http://web.utk.edu/~kcook/319S02/ex p3m.pdf (accessed September 30, 2015) [4] The University of Texas at Austin. Coordination Chemistry . http://www.geo.utexas.edu/courses/37 6m/coord_chem.htm (accessed September 30, 2015) [5] UC Davis ChemWiki. EDTA. http://chemwiki.ucdavis.edu/Inorganic_ Chemistry/Coordination_Chemistry/Lig ands/EDTA (accessed September 30, 2015)

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